J. Chem. Eng. Data 1980, 25,
+ +
are only a few vapor-liquid (VLE) data for the systems MCH toluene and MCH ethylbenzene. The VLE data for MCH benzene consist of isothermic measurements. There is a complete lack of (either isobaric or isothermal) VLE data for MCH o-xylene, MCH p-xylene, and MCH propylbenzene. A direct comparison of our results can be made with data of Funk et al. (2)for the MCH ethylbenzene system and less directly, because of differences in temperature, with the data of Schneider (5)for the MCH toluene system. To give a comparison with data obtained at different temperatures, we plotted Q0,5,the equimolar value of gEIRTvs. 1/Tin Figure 1, giving a straight line in the temperature interval considered. The slope of this line for the MCH toluene system gives HE,,, = 507 J. This average value in the 313-363 K interval agrees well who obtained = 523 J at with the data of Wdjcicki (9), a temperature of 298 K. For the second system the difference between our and the data of ~ ~ ~ (2) k is - about ~ t5 J. I t is evident from Figure 1 that such a difference lies within
+
+
+
+
+
+
+
161
161-163
the limits of experimental error.
Literature Cited Forziatti. A. F.. Norris. W. R.. Rossini. F. D., J . Res. Natl. Bur. Stand.. 43, 555 (1949). Funk, E. W., Chai, F. C., Prausnitz, J. M.. J. Chem. Eng. Data, 17, 24 (1972). Janaszewski, B., Oracz, P., Warycha, S.,5th Congress CHISA'75 Praha, 1975. ReDOrt NO. 3.3. Pitzer, K.'S., Curl, R. F., J . Am. Chem. SOC.,79, 2369 (1957). Schneider, G., Z.Phys. Chem. (Frankfurt am Main), 27, 171 (1961). Warycha, S., J . Chem. Thermodyn., 9, 813-820 (1977). Warycha, S., Z.Phys. Chsm. (Leipze),258, 864-872, 873-878 (1977). Willingham, C. J.. Taylor, W. J., Pignocco, F. D., J . Natl. Bur. Stand., 35, 219 (1945). WBjcicki, K., Sadowska, W., Bull. Acad, Pol. Sci., Ser. Sci. Chim., 18, 329 (1966).
Received for review August 16, 1979. Accepted January 10,1980. This work was supported by the Polish Academy of Sciences within the project PAN 03.10.
Ammonium Iodide-Potassium Iodide Mixed Crystals in Equilibrium with Aqueous and Aqueous-2-Propanol Solutions R. Ronald Richards,* Steven L. Young, Elizabeth Gebert, Felix Schreiner, and Evan H. Appelman Chemistry Division, Argonne National Laboratory, Argonne, Illinois 60439
+
+
The equilibria in the systems K I HN41 water and K I HN.1 equimolar mixtures of 2-propanol and water have been studied at 25 OC. The two iodides are completely miscible In the solid phase.
+
+
Introduction
In a recently proposed cycle for the thermochemical production of hydrogen ( 1, 2), the water-binding step is the reaction KI
+ H20 + NH, + COP
-
KHC03
+ NH41
This reaction is carried out in an equimolar mixture of water and 2-propanol, potassium bicarbonate precipitates, and the supernatant solution contains a mixture of K I and NH41from which the latter must be separated. The present investigation has been undertaken to elucidate the phase behavior of solid KI and NH41 in equilibrium with such solutions, as well as with purely aqueous solutions of the two salts. We hoped to find a means of precipitating NH41from the solutions that would be analogous to the method employed to recover NH,CI in the modified Solvay process (3). A study of the NH,CI-KCI-H,O and NH4Br-KBr-H20 systems has been reported ( 4 ) . I n both of these cases the potassium and ammonium salts exhibit limited mutual solubility. No comparable investigation of the iodide system appears to have been carried out.
Experimental Section Analytical reagent grade chemicals were used without further purification. Due to the limited mobility of ions within a solid and the different ratios of Kf/NH4+ in the solid and solvent phases, esAuthor to whom correspondence shoukl be addressed at the Department of Chemistry, Greenville College, Greenville, Illinois 62246
tablishment of equilibrium on mixing solid and solvent may be slow. Therefore equilibrium was approached by precipitation of solid from a solvent phase whose composition varied little during the experiment. Experiments with Equimolar Water -2- Propanol Solutions. Two burets were used to deliver simultaneously over 10 s 3 mL of 2-propanol and about 1.1 mL of aqueous KI-NH41 solution to a stirred flask thermostated at 25 OC. Volumes were such as to produce a final water/2-propanol mole ratio of 1:l. Precipitation of NH,I-KI occurred due to reduced solubility in the mixed solvent. The solution and solid were allowed to equilibrate for 30-60 s before the cycle was repeated. Since momentary deviations from equimolar 2-propanol-water on addition and mixing, with resultant inhomogeneity of the solid, are more likely when the solution volume is small, the solid was discarded after addition of the first 24 mL of 2-propanol. Cycles of simultaneous aliquot addition to the retained solution followed by equilibration were continued until sufficient precipitate had collected. The resulting mixture was stirred for 20-24 h, the solution was decanted, the soli was dried, and both phases were analyzed for NH,' K', and I-. The solvent mass loss on drying of the solid was used in conjunction with the density and composition of the mother liquor to correct the composition of the solid for deposition of salts from adhering mother liquor. Experiments with Aqueous Solutions. A stream of nitrogen was introduced over the surface of a stirred saturated aqueous solution of NH41 K I at 25 OC to induce evaporation of water and precipitation. Since solid that formed on the sides of the flask through evaporation would tend to have the composition of evaporated solution, care was taken to avoid contamination of the equilibrium solid phase distributed throughout the liquid with this solid. After the desired amount of precipitate was obtained, the nitrogen stream was stopped, the flask stoppered, and the mixture stirred for 3-20 h. One aliquot of liquid was removed for analysis on initiation of precipitation, and another was taken after the equilibration period. The results were averaged. Since the precipitated solid contained a higher pro-
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0021-9568/80/1725-0161$01.00/00 1980 American Chemical Society
162 Journal of Chemical and Engineering Data, Vol. 25, No. 2, 1980
+ Solvent
Table I. Solid-Solvent Equilibria for KI t ",I 2-propanol-water" [K']/[ 1-Ib solid
liquid
0.527 0.612 0.619 0.747 0.748 0.826 0.892 0.948 0.959 0.980
water [KC]/[ 1-lb solid
0.147 0.179 0.185 0.260 0.268 0.348 0.457 0.654 0.701 0.827
unit cell length, A
liquid
0.027 1 0.0595 0.103 0.149 0.189 0.285 0.437 0.585 0.731 0.847 0.908 0.930
0.0269 0.0535 0.101 0.137 0.151 0.247 0.365 0.493 0.624 0.75 1 0.830 0.870
Mixed Crystals as a
Table 11. Unit Cell Lengths of ",I-KI Function of Composition 7.258 7.225 7.181 7.144 7.102 7.060
[K+1 /[I- 1 0 0.185 0.398 0.594 0.802 1.00
(1)" (1) (4) (1) (1) (2)
"
Uncertainties are standard deviations. mole ratios. 4
I
,
Compositions are
I
a 50 mol % 2-propanol in water. b Compositions are mole ratios. IC
I
I
I
I
a,tAl 08-
-2 0 6 --> I
--. 1-
+504-
Flgure 2. Dependence of the cubic unit cell length ( a o )on the composition of mixed crystals of potassium and ammonium iodides. Compositions are mole ratios.
0 2-
0
02
I
1
I
04
06
08
IC
portion of Kf than the liquid, the second analysis was systematically higher in [NH4+]and lower in [K']; the average increase in [NH,'] was 0.8%. Analyses of the solid were corrected for the presence of mother liquor. Chemlcal Analysis. Iodide was determined by titration with standard silver nitrate using 0.5% Eosine Y indicator. The Kjeldahl method was used for ammonium ion. Potassium was determined either by atomic absorption spectroscopy or by addition of excess tetraphenyl boron and back-titration with a quaternary ammonium bromide solution in the presence of Clayton Yellow indicator (5). X-ray Analysis. Cell dimensions were obtained from Debye-Scherrer powder patterns taken with Cu Kcr radiation, using a Straumanis mounting. Results and Dlscusslon Composition data for the 2-propanol-water and water systems are given in Table I. The data are shown graphically in Figure 1. The KI-",I-H,O system exhibits one degree of freedom over the entire composition range at 25 O C and atmospheric pressure. For any given concentration of K+ in the solid, the Kf concentration in the liquid is fixed. The phase rule predicts two coexistent phases in a three-component system (KI, ",I, and solvent of fixed composition)with one degree of freedom. One solvent phase is present, so K I and NH41must form a solid solution at all compositions. Since the thermodynamic stability of the solid phase is independent of solvent (assuming no incorporation of solvent into the solid), the 2-propanol-water system verifies the conclusion over the composition range
studied. Continuous solid solution formation is also indicated by X-ray powder patterns of the solid. The unit cell lengths of the cubic structures are shown in Table I1 and Figure 2. The dimensions for the s o l i solutions are intermediate between those of pure KI and pure NH,I (6). The cell dimensions that we have obtained for pure NH,I and K I are respectively 0.014 and 0.073% lower than those of the literature values (6). This discrepancy is negligible, however, compared to the ca. 1% uncertainty in [K']/[I-] for the solid solutions. The complete miscibility of K I and NH,I in the solid phase contrasts with the behavior of the chloride and bromide systems, which exhibit limited mutual solubili, resulting in two soli phases when the solubili of one component in the other is exceeded (4). Since K I and NH41exist in the same crystalline form at 25 O C , while the chloride and bromide pairs do not (6), greater mutual solubili is expected in the i d i e system. The complete miscibility and identical crystalline forms of K I and NH41are probably due to the greater interstitial volume created by iodide ions as compared with the smaller chloride and bromide ions. With regard to the original motivation of the study, our results indicate that it will not be possible to isolate NH41by selective precipitation from solutions containing both NH,I and KI. Hence the cycle for thermochemical hydrogen generation to which we referred earlier will have to rely on other means of separation. Acknowledgment We wish to thank the Argonne Analytical Group for the ammonium ion analyses and for the potassium ion analyses by atomic absorption. R. R. Richards and S. L. Young acknowledge the support of the Argonne Center for Educational Affairs. Literature Clted (1) Schreiner, F., Appelman, E. H., Sum. f r o g . Chern., 8, 171 (1977). (2) Appeiman, E. H., Schreiner, F., Abraham, B. M., "Proceedings of the 2nd Workl Hydrogen Energy Conference", Zurich, Switzerland, 1978, T. N. Verziroglu and W. Seifritz, Ed., Pergamon Press, New York, 1978, pp 851-907.
J.
(3) Kirk-Othmer. “EncvcloDedia of Chemical Technoloav”. -. 2nd ed.. 1963, pp 734-40. (4) Flatt, R., Burkhardt, G., Helv. Chirn. Acta, 27, 1605 (1944). (5) “Official Methods of Analysis”, 10th ed., Association of Official Agricultural Chemists, Washington, D.C., 1965, p 23. Shall, E. D., Anal. Chem., 29, 1044 (1957). I
,
I
163
Chem. Eng. Data 1980, 25, 163-170
.
(6) Merrill, L., J . Phys. Chern. Ref. Data 8, 1205 (1977). Received for review October 2, 1979. Accepted February 4, 1960. This research was performed under the auspices of the Office of Basic Energy Sciences, Division of Chemical Sciences, U.S. Department of Energy.
+
Vapor-Liquid Equilibrium for the Binary Systems Propionitrile Ethylbenzene and Acetonitrile 4- Ethyl Acetate, Ethyl Alcohol, Toluene and
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01 Muthu, Patrick J. Maher, and Buford D. Smith’ Thermodynamics Research Laboratory, Washington University, St. Louis. Missouri 63 730
Total pressure vapor-liquid equilibrium data were measured for the following four binaries: ethyl acetate acetonltrile at 313.15, 353.15, and 393.15 K; ethyl alcohol acetonitrile at 293.15, 343.15, and 393.15 K; acetonitrile toluene at 293.15, 343.15, and 393.15 K; proplonltrlle ethylbenzene at 313.15, 353.15, and 393.15 K. The P , T , x data were reduced to y , y,and GE values by the Mlxon-Gumowskl-Carpenter method. The vlrlal equation of state truncated after the second coefflclent was used to calculate the vapor-phase fugacity coefficients. The Tsonopoulos correlation was used to predict the second vlrlal coefflcients. For comparison purposes, the data for the acetonltrlle toluene system were also reduced by the Barker method using a four-suffix Margules equation.
+
+
+ +
+
Table I. Chemicals Used component
vendor
stated purity, %
acetonitrile ethyl acetate ethyl alcohol ethylbenzene propionitrile toluene
Burdick and Jackson Burdick and Jackson U.S. Industrial Chemicals El Paso Products Co. Aldrich Burdick and Jackson
99.9+ 99.9 200 proof 99.96 99.0 99.9
uncertainties for the three reported variables ( P , T, and x ) due to resolution limits in the instrumentation were summed to give an expression for the total effect on the reported vapor pressures. The total possible resolution error in kPa is given by total possible resolution error = 0.03
Il +
0.0005 -
Introduction The selection of the systems covered in this paper was based on the group interaction parameter grid of the UNIFAC correlation scheme outlined by Fredenslund et al. ( 7). The aim was to fill in places in the grid where either sparse or no data were available. Data were measured at three temperatures for each of the four systems. Maher and Smith (2,3) have reported similar data previously for binary systems containing chlorobenzene and aniline.
Experimental Apparatus and Techniques The apparatus used for this study has been described previously ( 4 ) . Fifteen discrete cells are loaded with the two pure components and thirteen intermediate binary mixtures. The mass of each component loaded in a cell is determined by the use of an analytical balance. The cells are attached to a low-volume manifold, and the contents are degassed in situ by successive freezing-evacuation-thawing cycles. When the degassing is completed, the manifold assembly is placed in a constant-temperature bath. Pressure measurements are made by opening each cell in turn to a differential pressure transducer used as a nulling device. The nitrogen pressure used to balance the cell vapor pressure is measured by using a separate transducer that is calibrated frequently vs. a dead weight gauge. The bath temperature is measured with a platinum resistance thermometer that is calibrated vs. a local platinum resistance temperature standard using the IPTS-68 temperature scale. The uncertainties in the measured temperatures and mole fractions are f0.03 K and f0.0005,respectively. The estimated
0.00035P+ 0.003
+
6.89476[(2.5)(resolution for 572)]
where P = experimental pressure, (aP/a T), = temperature derivative of pressure at constant composition, and (aP/ax), = absolute value of the composition derivative of pressure at constant temperature. The first two terms provide the uncertainty in the pressure measurement due to the possible temperature and mole fraction errors. The last three terms are related to the uncertainty in the measurement of pressure with a Datametrics 572 transducer. That transducer reads out psia and the “resolution” at various pressure levels is as follows: 0.0001 below 4 psia, 0.001 below 40 psia, 0.01 below 400 psia, etc. The error in the reported vapor pressures, as computed by the above equation, will be referred to as the “Resolution Error Band” (REB) to distinguish it from the total possible experimental error which will include other factors such as residual gas, residual water, and chemical reaction. The width of the REB is usually about f0.1 YO or less of the system pressure.
Chemicals Used Table I lists the chemicals used and their stated purity. All chemicals except propionitrile were available in at least 99.9 YO purity. Activated Linde molecular sieves (either 3A or 4A) were put in the containers with the chemicals as they were received. Prior to being loaded in the metal cells, the chemicals were distilled over freshly activated molecular sieves through a Vigreaux column (25 mm 0.d. and 470 mm long) with the first and last portions being discarded. The distilled samples were
0021-9568/80/1725-0163$01.00/00 1 9 8 0 American Chemical Society
